The Covalent Bonding Model Lecture Notes

Lewis Structures and Valence Electrons

• The Lewis structure of an atom represents its valence electrons, which occupy the outermost energy level and are responsible for chemical bonding.
• Atoms lose, gain, or share electrons to attain a noble gas configuration.
• The Octet Rule: Most atoms seek to have $8$ electrons in their outermost energy level, except for hydrogen and helium.
• Lewis structures for the first 20 elements (e.g., $H\bullet$, $He:$, $Li\bullet$, $Be:$, etc.) follow periodic group trends.

The Covalent Bonding Model

• Covalent bonding involves a shared pair of electrons between two atoms and is the predominant bond in molecules.
• An atom contributes the same number of electrons for sharing as it needs to complete its octet.
• Bond Types:
  • Single bond: $1$ pair of shared electrons.
  • Double bond: $2$ pairs of shared electrons.
  • Triple bond: $3$ pairs of shared electrons.

Procedures for Writing Lewis Structures

• Step 1: Calculate total valence electrons. Add electrons for negative charges; subtract for positive charges.
• Step 2: Write the skeletal arrangement. Choose the "leftist or lowest" element as the central atom. Connect atoms with single bonds.
  • Hydrogen forms only one bond.
  • Oxygen usually forms a maximum of two bonds.
• Step 3: Subtract used electrons from the total. Distribute remaining electrons as lone pairs to terminal atoms until they have an octet, then place any leftovers on the central atom.

Exceptions and Resonance

• Expanded Valence Shell: Central atoms in Period 3 or below can use d orbitals to hold $10$ or $12$ electrons.
• Electron Deficient/Free Radicals: Highly reactive compounds with fewer than $8$ electrons or odd numbers of electrons.
• Resonance: The movement of electrons that stabilizes a molecule (e.g., $O_3$).
• Polyatomic Ions: Groups of atoms with a charge that behave as a single unit (e.g., $NH_4^+$, $SO_3^{2-}$, $NO_2^-$, $NO_3^-$, $I_3^-$).

Electronegativity and Bond Polarity

• Electronegativity (EN): The relative attraction an atom has for a shared pair of electrons. It generally increases from left to right across a period.
• Bond Types by EN Difference ($\Delta EN$):
  • Nonpolar Covalent: Equal sharing, $\Delta EN \leq 0.5$.
  • Polar Covalent: Unequal sharing, $0.5 < \Delta EN < 1.9$. Results in partial positive and negative charges.
  • Ionic: Minimal sharing/electron transfer, $\Delta EN > 2.0$ (or $1.9$ to $2.0$).
• Formal Charge (FC): $FC = \text{valence electrons} - (\text{electrons} + \text{bonds})$.

Molecular Geometry

• $H_2O$: V-shaped.
• $CO_2$: linear shape.
• $BF_3$: trigonal planar.
• $CH_4$: tetrahedral.

• Lewis Structures and Valence Electrons

  • The Lewis structure of an atom represents its valence electrons, which occupy the outermost energy level and are responsible for chemical bonding.
  • Atoms lose, gain, or share electrons to attain a noble gas configuration.
  • The Octet Rule: Most atoms seek to have $8$ electrons in their outermost energy level, except for hydrogen and helium.
  • Lewis structures for the first 20 elements (e.g., $H\bullet$, $He:$, $Li\bullet$, $Be:$, etc.) follow periodic group trends.

• The Covalent Bonding Model

  • Covalent bonding involves a shared pair of electrons between two atoms and is the predominant bond in molecules.
  • An atom contributes the same number of electrons for sharing as it needs to complete its octet.
  • Bond Types:
  • Single bond: $1$ pair of shared electrons.
  • Double bond: $2$ pairs of shared electrons.
  • Triple bond: $3$ pairs of shared electrons.

• Procedures for Writing Lewis Structures

  • Step 1: Calculate total valence electrons. Add electrons for negative charges; subtract for positive charges.
  • Step 2: Write the skeletal arrangement. Choose the "leftist or lowest" element as the central atom. Connect atoms with single bonds.
  • Hydrogen forms only one bond.
  • Oxygen usually forms a maximum of two bonds.
  • Step 3: Subtract used electrons from the total. Distribute remaining electrons as lone pairs to terminal atoms until they have an octet, then place any leftovers on the central atom.

• Exceptions and Resonance

  • Expanded Valence Shell: Central atoms in Period 3 or below can use d orbitals to hold $10$ or $12$ electrons.
  • Electron Deficient/Free Radicals: Highly reactive compounds with fewer than $8$ electrons or odd numbers of electrons.
  • Resonance: The movement of electrons that stabilizes a molecule (e.g., $O_3$).
  • Polyatomic Ions: Groups of atoms with a charge that behave as a single unit (e.g., $NH_4^+$, $SO_3^{2-}$, $NO_2^-$, $NO_3^-$, $I_3^-$).

• Electronegativity and Bond Polarity

  • Electronegativity (EN): The relative attraction an atom has for a shared pair of electrons. It generally increases from left to right across a period.
  • Bond Types by EN Difference ($riangle EN$):
  • Nonpolar Covalent: Equal sharing, riangle EN ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } riangle EN ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ } ext{ }
  • Polar Covalent: Unequal sharing, 0.5 < riangle EN < 2.0.
  • Ionic: Minimal sharing/electron transfer, riangle EN > 2.0$(or$1.9$to$2.0).

• Formal Charge (FC): FC = ext{valence electrons} - ( ext{electrons} + ext{bonds})$$.